Chemical properties of s-block elements

1) All the alkali metals are highly reactive elements since have a strong tendency to lose the single valence s-electron to form unipositive ions having inert gas configuration.

2) This reactivity arises due to their low ionisation enthalpies and high negative values of their standard electrode potential.

3) Due to their strong tendency to lose the single valence electron ,alkali metals also act as a strong reducing agents. Reactivity of an element is measured in terms of its reducing character.

4) Ionization enthalpy is a measure of the tendency of an atom to lose electrons in the gaseous state. Thus, lower the ionization enthalpy, greater is the tendency of an element to lose electron and hence stronger is the reducing character or higher is the reactivity of the element.

5)The ionization enthalpies of alkali metal decreases down the group, therefore, their reducing character or reactivity in the gaseous state increases from Li to Cs , i.e. Li< Na< K< Rb<Cs.

Li is the strongest while sodium is the least powerful reducing agent in aqueous solution.

Electrode potential is a measure of the tendency of an element to lose electrons in the aqueous solution. More negative is the electrode potential, higher is the tendency of the element to lose electrons and hence stronger is the reducing agent.

Since the standard electrode potential of alkali metals become more and more negative as we move down the group from Na to Cs, therefore ,reducing character of these elements increases in the same order i.e. Na to Cs.

Standard electrode potential of Li, is the lowest i.e. i.e. -3.04 V.

Lithium is the strongest reducing agent in the aqueous solution.

 

Electrode potential depends upon

1) enthalpy of sublimation

2) ionization enthalpy and

3) enthalpy of hydration.

Li(s) ———-> Li(g)

Li(g) —–> Li⁺ (g) + e‾

Li⁺ (g) +aq —–> Li⁺ (aq) + e‾

The sublimation enthalpies of alkali metals are almost similar. Since Li has the smallest ionic size among alkali metal ,its enthalpy of hydration is the highest. Although ionization enthalpy of Lithium is the highest among alkali metal, it is more than compensated by the large hydration enthalpy released.

Among alkali metals, Lithium has the most negative standard electrode potential and hence is the strongest reducing agent in the aqueous solution. Stronger reducing agent have higher reactivity ,therefore, Lithium should also be most reactive in aqueous solution.

Reactivity towards water

Alkali metals react with water, liberating dihydrogen and forming their hydroxides. The reaction becomes more and more violent as we move down the group. Lithium reacts gently ,sodium melts on the surface of water and the molten metal moves around vigorously and may sometimes catch fire ,potassium melts and always catches fire and so are Rb and Cs.

2Li + 2H₂O ——> 2LiOH + H₂

2 Na + 2H₂O ——> 2NaOH + H₂

2K + 2H₂O ——> 2KOH + H₂

 

Because of their large negative reduction potential ,alkali metals are better reducing agents than hydrogen. They react with compounds containing acidic hydrogen atom such as water ,alcohol and acetylene liberating hydrogen gas.

2M + 2H₂O ——> 2MOH + H₂

2M + 2C₂H₅OH ——>2 C₂H₅OH + H₂

2M + HC≡CH ——-> MC≡CM + H

Lithium is the least reactive while the reactivity of other alkali metal towards water and other acidic hydrogen containing compounds increases on moving down the group from Na to Cs.

Standard electrode potential (E°) and Gibbs free energy (ΔG°) are related by the equation,

ΔG° = – nFE°

Where n is the number of electrons lost by the metal and F is the Faraday constant.

Since E° for the reaction, Li⁺ (aq) + e‾ ———> Li(s)  has the lowest negative value, i.e. -3.04 V , therefore, ΔG° of the reaction has the largest positive value. This reaction does not occur.

The reverse reaction , Li (s) ———> Li⁺ (aq) + e‾ , has a large negative value of ΔG°, so Li liberates more energy than any other alkali metal when it reacts with water.

Que Why Li reacts with water gently ,whereas Na which liberates less energy, reacts more vigorously and the hydrogen produced catches fire.

Ans Sodium has a low melting point, and the heat of the reaction is sufficient to make it melt or even vaporise. The molten metal thus spread out thereby exposing a larger surface to water. As a result, it reacts even faster, gets even hotter and catches fire. Lithium has high melting point. Although the heat of the reaction is high, it is still not sufficient to melt the metal and hence the reaction proceeds gently.

Reactivity towards oxygen

The alkali metals tarnish in air due to the formation of an oxide or hydroxide on the surface. When heated in excess of air ,alkali metals form different types of oxides depending upon the nature of the metal :

1) Lithium when heated in oxygen form Lithium monoxide

4Li + O₂ —–> 2 Li₂O

2) Sodium when heated with oxygen at about 575 K forms mainly  sodium peroxide

2 Na + O₂ —–> 2 Na₂O₂

Other alkali metals i.e. K, Rb, Cs react with oxygen to form superoxides of the formula MO₂ where M= K, Rb, Cs.

The reactivity of alkali metals with oxygen increases down the group. The increasing stability of peroxides or superoxide, as the size of the metal cation increases ,is due to the stabilisation of larger anion by larger cations through higher lattice enthalpy.

Reason: Because of the small size , Li⁺ as a strong positive around it which attracts the negative charge so strongly that it does not permit the oxide anion O^2- to combine with another oxygen atom to form peroxide ion, O₂^2- . Na⁺ ion because of its large size than Li⁺ ion has weaker positive field around it which cannot prevent O^2- ion to combine with another oxygen atom to form peroxide ion O₂^2-

 

The larger K⁺ , Rb⁺ and Cs⁺ ions have still weaker positive fields around them which cannot prevent even peroxide ions, O₂^2- to combine with another oxygen atom to form superoxide.

The superoxide ion has a 3 electron bond i.e. it has one unpaired electron which makes it colour and paramagnetic. The normal oxides of alkali metals are colourless and diamagnetic.

All alkali metals except Li form ozonides of the formula , MO₃

M(s) + O₃ (g) ——> MO₃

These ozonides are unstable and decompose on standing to form superoxide and oxygen

2 MO₃(s)———–> 2 MO₂(s) + O₂ (g)

Reactivity towards air and moisture

All the alkali metals on exposure to atmosphere gets converted into oxides, hydroxide and finally to carbonates. Thus  alkali metals get tarnished when exposed to air and moisture.

4 M + O₂ ——> 2 M₂O

M₂O + H₂O —–> 2 MOH

2 MOH + CO₂ —–> M₂CO₃ + H₂O

Alkali metals are stored in inert hydrocarbon solvents like petroleum, ether, hexane benzene and kerosene oil which prevents them from coming in contact with air and moisture.

Reactivity towards hydrogen

All the alkali metals react with hydrogen at 673 K to form colourless crystalline ionic hydrides of the general formula MH Where M stands for the alkali metals.

2 M + H₂ ——-> 2 MH

where M = Li, Na, K, Rb , Cs

1) The order of reactivity of the alkali metals towards hydrogen decreases as we move down the group from Li to Cs .The lattice enthalpy of these hydrides decreases progressively as the size of the metal cations increases and thus the stability of these hydrides decreases from LiH to CsH.

2Li + H₂ ——> 2 LiH

2 Na + H₂ ——> 2 NaH

2) All the alkali metal hydrides are ionic solids with high melting points.

The ionic character of the hydrides ,however ,decreases from Li to Cs.As the size of the cation increases, the anion hydride ion can polarize the cation more easily.As a result, the covalent character increases and hence the ionic character deceases.

3) All the hydrides behave as strong reducing agents and their reducing nature increases down the group.

4) Since these hydrides contain the hydride ion ,therefore ,they liberate hydrogen at the anode on electrolysis.

5) All these hydrides react with Proton Donor such as water ,alcohol ,gaseous ammonia and alkynes liberating H₂ gas.

LiH (s) + H₂O (l) ——> LiOH (aq) +  H₂ (g)

NaH(s) + ROH (l) ——–> RONa (s) +H₂ (g)

NaH(s) + NH₃ (g) ——> NaNH₂ (s) + H₂ (g)

2 KH (s) + HC≡CH (g) ——–> KC≡CK (s) + 2 H₂ (g)

Lithium hydride is used as a source of hydrogen for military purpose and for filling metrological balloons since it has a low molecular weight and on reacting with water ,it evolves highest percentage of hydrogen by weight.

 

Reactivity towards halogens

Alkali metals react vigorously with halogens to form ionic metal halides of the general Formula , MX where M stands for metal and X for the halogens.

2 M + X₂ ———> 2 MX

The reactivity of alkali metal towards a particular halogen increases as we move down the group from Li to Cs due to a decrease in the ionization enthalpy or increase in the electropositive character of the metal. Potassium reacts with chlorine more vigorously than sodium.

Reactivity of halogens towards a particular alkali metal decreases from F₂ to I₂ i.e.F₂ > Cl₂ > Br₂ > I₂.

Solution in liquid Ammonia

All the alkali metal dissolve in liquid ammonia giving highly conductive deep blue solutions.

These solution contain ammoniated cations and ammoniated electrons.

M + (x+y)NH₃ ——> M⁺(NH₃)_x + e‾(NH₃)_y

When ordinary light fall on these ammoniated electrons, they get excited to higher energy level by absorbing energy corresponding to red region of the visible light. As a result ,transmitted light is blue which imparts blue colour to the solution.

1) Dilute solution of alkali metals in liquid ammonia are dark blue in colour but as the concentration increases above 3 M ,the colour changes to copper bronze and the solution acquires metallic lustre due to formation of metal ion clusters.

2) The blue coloured solutions are paramagnetic due to the presence of large number of unpaired electrons but bronze solutions are diamagnetic due to formation of electrons cluster in which ammoniated electrons with opposite Spin group together.

2 e‾ (NH₃) _y ———-> [ ↑ e‾ (NH₃) _y ] [ ↓ e‾ (NH₃) _y ]

3) The solution of alkali metals in liquid ammonia are good conductor of electricity due to the presence of ammoniated cations and ammoniated electrons. The conductivity decreases as the concentration increases since the ammoniated metal cations and bound by the free unpaired electrons which have been described as expanded metals.

4) The solution of alkali metals in liquid ammonia are stronger reducing agents than hydrogen and hence will react with water to liberate hydrogen.

5) In the presence of impurities or catalyst such as iron ,the blue coloured solution decompose to form metal amides with the liberation of dihydrogen.

6) Under anhydrous condition and in the absence of catalytic impurities, these solutions are stable and can be stored for several days.